(Original post by Miss Understood)
Thanks, just to clarify a few things:

Why is Carbon-12 the sole measurement of relative atomic mass and not any other element?

How did scientist work out the relative atomic mass of every element and how did that soon correlate to working out the quantities of protons/neutrons in an atom?

Other elements have been used in the past. I know I shouldn't use WIKI as a source but it is getting late...

The atomic weight scale has traditionally been a relative scale, that is without an explicit unit, with the first atomic weight basis suggested by John Dalton in 1803 as 1H.[3] Despite the initial mass of 1H being used as the natural unit for atomic weight, it was suggested by Wilhelm Ostwald that atomic weights would be best expressed in terms in units of 1/16 weight of oxygen. This evaluation was made prior to the discovery of the existence of elemental isotopes, which occurred in 1912.[3]

The discovery of isotopic oxygen in 1929 led to a divergence in atomic weight representation, with isotopically weighted oxygen (chemistry) and pure 16O (physics) bases both used as the basis for the atomic mass unit (amu). The inevitable divergence could result in errors in computations, and was thus unwieldy. The reference was changed to carbon-12 in 1961[4] and a new symbol "u" replaced the now deprecated "amu".

1 u = mu = 1/12 m(12C)

The current unit is referred to as the "unified atomic mass unit" u.[5] The choice of carbon-12 was used to minimise further divergence with prior literature.[3]

(Original post by Miss Understood)
Thanks, just to clarify a few things:

Why is Carbon-12 the sole measurement of relative atomic mass and not any other element?

How did scientist work out the relative atomic mass of every element and how did that soon correlate to working out the quantities of protons/neutrons in an atom?

I don't think there's anything magical about the choice of Carbon-12, just an accident of history. Before adopting Carbon-12, 1/16th the mass of oxygen was used. Unfortunately that standard was adopted before isotopes were fully understood, so the unit was defined differently depending on whether you considered pure Oxygen-16 or the natural mixture of oxygen isotopes. Carbon-12 was eventually decided on as a new standard to replace the confusion. I'm not sure exactly why although I think I think basing the atomic mass unit on 1/12 of Carbon-12 gives a similar answer to basing it on 1/16 the mass of the naturally occurring mix of Oxygen isotopes.

Google 'mass spectrometry' if you want to read about how atomic masses are usually measured. I'm not quite sure I understand your last question. Because of the way the atomic mass unit (u) is defined, nucleons (protons and neutrons) have an approximate mass of 1u. So if you know the atomic mass of a certain isotope of an atom, you can work out how many protons and neutrons it has (for example He-4 has a relative atomic mass of 4.003, so we know there must be a total of 4 nucleons).

The method may not work when dealing with "relative atomic weight" as this usually refers the average mass (in atomic mass units) of the various isotopes (weighted according to their natural abundance). It is these values which tend to appear on periodic tables. The wikipedia articles on atomic mass and atomic weight are a good place to start if you want to know more, as someone has said.

(Original post by js374)
You win the award for clearest post of this thread. If you will forgive me for saying it, I think you may have got confused because the initial question you asked was a little ambiguous so the answers were fairly divergent.

The phrase relative is used because all of the masses used are 'relative' to the mass of carbon-12. The exact mass of an atom of carbon-12 would be ~1.396x10^-17 g. This is clearly a very tiny number and is a pain to work with so it is much simpler to use the relative system. Take water as an example, the relative mass is ~18* so this is ~18* times heavier than 1/12 of a carbon atom. The weight of everything is calculated relative to the carbon atom weight. Exactly 12 just means that there are no decimal places and it has not been rounded.

Relative atomic mass take into account isotopic differences. So while carbon-12 has a RAM of 12, carbon does not and has a RAM of 12.011 to take account of the isotopes.

How scientists worked out the numbers was a long programme of research by different scientists. Lavoisier, Dalton, Avagadro and Mendeleev are some of the more important people if you want to read up on it.

I hope that is clearer.

* I use "~" because it is not exactly 18 and but i can't be bothered to work it out/look it up but I also don't want Charco or Phil jumping down my throat about how it is not a whole number

You've really cleared this up a bit.

I just have to ask 2 more stupid questions:
In carbon 12, where did the value 12 atomic mass unit come from?

So by definition carbon 12 isn’t a RAM, it’s just an atom. So they do scientists say it is?

The masses quoted in the periodic table are relative atomic masses in grams per mole. One mole of a given substance contains ~6*10^23 particles; this is Avagadro's constant. Therefore, if you want to know the mass of one atom of a given element, you multiply that element's relative atomic mass by the reciprocal of Avagadro's constant. This will give you the mass of one atom of a given element in grams. Then, to move into SI units, we must further multiply by 10^(-3) to give the mass of one atom of a given element in kilograms.

Hydrogen has a relative atomic mass of roughly 1. I believe this means that one mole of hydrogen atoms has a mass roughly equal to 1/12 of the mass of one mole of carbon-12 atoms at rest. It also means that one mole of hydrogen atoms has a mass of roughly one gram. So to find the mass of just a single hydrogen atom:

To try and justify this to yourself, pay special attention to the units involved. Our first quantity is grams per mole. We then divide by some number of atoms per mole, which is the same as multiplying by some other number of moles per atom. The "per mole" of the first term and "moles" of the second term cancel out to give "grams per atom", or the mass of a single atom in grams. We divide by some number of grams per kilogram, which is the same as multiplying by some other number of kilograms per gram. The "grams" of the first two terms cancel with the "per gram" of the third term, leaving "kilograms per atom", or the mass of just one atom in kilograms. This sort of dimensional analysis is often useful when converting between different units.

In answer to your second question, one method of working out relative atomic masses is actually by using mass spectroscopy. If we define one atomic mass unit (that is, one unit of relative atomic mass, Ar) to be one twelfth the mass of a carbon-12 atom (noting that is is not a contradiction to what I said in the first paragraph; if one mole of a substance of Ar 1 has mass equivalent to 1/12 that of one mole of a substance of Ar 12, then one atom of a substance of Ar 1 too has mass equivalent to 1/12 that of one atom of a substance of Ar 12), then by knowing the charge of an ion passing through a mass spectrometer (for your example we would have a cation of Cl-37, charge 1+) and measuring how far it is deflected by an electromagnetic field of some known magnitude and orientation, we can work backwards mathematically to find the mass-charge ratio (m/z) which in this case would actually be equal to the mass of Cl-37, given that z=1.

Finally, both 37 and 35.5 are relative atomic masses. They are the Ar for different things, however. 37 refers to the relative atomic mass of chlorine-37 only, whereas 35.5 is the weighted average (based on relative abundances in nature) of all known isotopes of chlorine. If you were dealing with a specially prepared sample of pure chlorine-37, your calculations would take Ar 37 (although 37 isn't the exact number, but rather a good approximation). If, on the other hand, you were doing calculations based on, say, the chlorine dissolved in a swimming pool, you'd go for Ar 35.5.

I hope that has helped to clear up some confusion. I'm only in my first year of chemistry, so if anyone more experienced (or less for that matter!) would care to point out any errors, I'd appreciate it.

Last edited by StandardCarpet; 04-04-2011 at 13:33.
Reason: Missed out a minus sign in the first paragraph.

(Original post by Miss Understood)
Wow thanks for this. Are you studying at AS level? I'm actually resitting my A-level chemistry and I think I'm look too much into everything. I seen stupid compared to you. You'll Ace your papers.

Haha, no way! If I knew all that stuff when I was sitting AS you're right, I would have aced my papers. I'm in my first year at uni, but that doesn't mean that I get it right every time - chemistry is a difficult subject!

It's no problem, I sometimes ask for help on here so when the rare opportunity crops up to repay the favour I really don't mind at all. And you can never look too much into anything; just remember that at A-level there are a number of "required" answers that you really just need to regurgitate from textbooks. Study away, but make the distinction between what you "need" to know and what's actually going on.

(Original post by StandardCarpet)
Haha, no way! If I knew all that stuff when I was sitting AS you're right, I would have aced my papers. I'm in my first year at uni, but that doesn't mean that I get it right every time - chemistry is a difficult subject!

It's no problem, I sometimes ask for help on here so when the rare opportunity crops up to repay the favour I really don't mind at all. And you can never look too much into anything; just remember that at A-level there are a number of "required" answers that you really just need to regurgitate from textbooks. Study away, but make the distinction between what you "need" to know and what's actually going on.

(Original post by Mbob)
Going off topic a bit..., so OP please ignore:

Unless I misunderstand you (which is possible, it's very late!), I don't think this is right. The rest masses of protons and neutrons in an atom don't add up the rest mass of the atom because of the mass-energy equivalence you talk about. Work is done in binding the nucleons together, i.e potential energy of the system is lost. This is exhibited as loss of mass, hence the mass of bound nucleons is less than the sum of the individual nucleons. It is not, I believe, due to rounding errors in the definition of the atomic mass unit.

Yea you guys got it spot on, I totally forgot about binding energies which would of course reduce the mass of carbon 12 thanks for spotting that for me